NSSCO Chemistry (6117) cheat sheet

Matter is anything with mass and volume; solids vibrate about fixed positions, liquids slide past each other, and gases move randomly at high speed.

A pure substance has fixed composition: an element contains one type of atom, while a compound contains two or more elements chemically bonded.

A mixture contains substances not chemically bonded; each component keeps its own properties and can be separated by physical methods.

A physical change alters state, size or shape without forming a new substance; a chemical change forms a new substance and usually involves an energy change.

Pure substances melt and boil at sharp fixed temperatures at a given pressure; impurities lower the melting point, broaden the melting range and raise the boiling point.

Sublimation is direct change from solid to gas without a liquid stage; iodine, ammonium chloride and dry ice sublime under suitable conditions.

Diffusion is net movement of particles from high to low concentration due to random motion; it is faster at higher temperature and for lower particle mass.

Density = mass / volume; common units are g/cm3 and kg/m3, with 1 g/cm3 = 1000 kg/m3.

On a heating curve, temperature stays constant during melting and boiling because added energy breaks forces between particles rather than raising kinetic energy.

Filtration separates an insoluble solid from a liquid: the solid is the residue on the filter paper and the liquid is the filtrate.

Evaporation removes solvent from a solution; crystallisation gives purer crystals by cooling a saturated solution slowly.

Simple distillation separates a solvent from a solution by boiling and condensing; fractional distillation separates miscible liquids with different boiling points.

A separating funnel separates immiscible liquids by density; the lower-density layer is run off after the stopcock is opened carefully.

Chromatography separates soluble substances by different rates of movement; Rf = distance travelled by spot / distance travelled by solvent front.

Decantation pours liquid off a settled solid; centrifugation speeds up settling of fine suspended particles.

A pure solid has a sharp melting point and a pure liquid has a sharp boiling point; chromatography can show whether a substance is pure.

Colourless chromatography spots may need a locating agent or ultraviolet light before their positions can be measured.

An atom has a nucleus containing protons and neutrons, surrounded by electrons; proton charge +1, neutron 0, electron -1.

Atomic number Z = number of protons = number of electrons in a neutral atom; it identifies the element.

Mass number A = protons + neutrons; neutrons = A - Z.

Isotopes are atoms of the same element with the same Z but different A; they have identical chemical properties but different masses.

Electron shells fill 2, 8, 8 for the first 20 elements; sodium is 2,8,1 and chlorine is 2,8,7.

The number of outer-shell electrons determines the group of a main-group element; the number of occupied shells determines the period.

An ion forms when an atom gains or loses electrons: cations are positive and anions are negative.

Relative atomic mass Ar is the weighted mean mass of an element's isotopes compared with 1/12 the mass of a carbon-12 atom.

Rutherford's alpha-scattering experiment showed that most alpha particles pass straight through but a few deflect strongly, indicating a small dense positive nucleus.

An ionic bond is the electrostatic attraction between oppositely charged ions formed by electron transfer, usually between a metal and a non-metal.

A covalent bond is a shared pair of electrons between atoms, usually between non-metal atoms.

A metallic bond is attraction between a lattice of positive metal ions and a sea of delocalised electrons.

Ionic compounds have high melting points, conduct when molten or aqueous, do not conduct when solid, and are often soluble in water.

Simple molecular substances have weak intermolecular forces, low melting and boiling points, and usually do not conduct electricity.

Giant covalent structures such as diamond and silicon dioxide have high melting points; graphite conducts because of delocalised electrons, while diamond does not.

Dot-and-cross diagrams show outer-shell electrons; ionic diagrams show charges and brackets, while covalent diagrams show shared electron pairs.

A polar covalent bond has unequal electron sharing due to an electronegativity difference; a symmetrical molecule can still be non-polar overall.

One mole is the amount containing 6.02 × 10^23 particles, the Avogadro constant.

n = m / M, where n is moles, m is mass in grams and M is molar mass in g/mol.

For a gas, n = V / molar gas volume; at room temperature and pressure the molar gas volume is often 24 dm3/mol, but use the value supplied in a question.

Concentration c = n / V, with V in dm3 and c in mol/dm3; therefore n = c × V.

Percentage yield = (actual mass of product / theoretical mass) × 100; atom economy = (Mr of desired product / total Mr of reactants) × 100.

Empirical formula is the simplest whole-number ratio of atoms; molecular formula = empirical formula × n, where n is found from Mr.

A balanced equation gives the mole ratio of reactants and products, allowing mass, volume or concentration calculations.

The limiting reactant is used up first and determines the maximum product; the other reactant is in excess.

Heating a hydrated salt drives off water of crystallisation; the mass loss gives the moles of water per mole of anhydrous salt.

An acid is a proton donor; a base is a proton acceptor; an alkali is a soluble base that produces hydroxide ions in water.

The pH scale is 0-14: pH below 7 is acidic, pH 7 is neutral and pH above 7 is alkaline.

Neutralisation is acid + base -> salt + water; the ionic equation is H+ + OH- -> H2O.

Acid + reactive metal -> salt + hydrogen; acid + carbonate -> salt + water + carbon dioxide.

A strong acid ionises completely in water; a weak acid ionises only partially, so at the same concentration the weak acid has the higher pH.

Soluble salts can be made by reacting an acid with excess insoluble base, carbonate or metal, then filtering and crystallising.

Soluble salts with two soluble reactants are made by titration: add acid to alkali to the endpoint, repeat without indicator, then crystallise.

Insoluble salts are made by precipitation, mixing solutions that contain the required ions.

Solubility rules: all nitrates are soluble; most chlorides are soluble except AgCl and PbCl2; most sulfates are soluble except BaSO4 and PbSO4; most carbonates are insoluble except sodium, potassium and ammonium carbonates.

Elements are arranged in order of increasing atomic number; vertical columns are groups and horizontal rows are periods.

For main-group elements, group number equals the number of outer-shell electrons; period number equals the number of occupied electron shells.

Group 1 alkali metals are soft, low-density, reactive metals; reactivity increases down the group and they form +1 ions.

Group 7 halogens are diatomic and reactive non-metals; reactivity decreases down the group.

A more reactive halogen displaces a less reactive halogen from its salt solution; chlorine displaces bromide, but iodine cannot displace chloride.

Group 0 noble gases have full outer shells, are monatomic and are very unreactive.

Transition metals have variable oxidation states, coloured compounds, catalytic activity and high density or melting point.

Across a period, metallic character decreases and non-metallic character increases; oxides change from basic through amphoteric to acidic.

Metalloids lie near the staircase line and have properties intermediate between metals and non-metals.

Reactivity series order: K, Na, Ca, Mg, Al, C, Zn, Fe, Sn, Pb, H, Cu, Ag, Au.

Metals above carbon in the reactivity series are extracted by electrolysis; metals below carbon are extracted by reduction with carbon.

Iron is extracted in a blast furnace: haematite Fe2O3 is reduced by carbon monoxide, and limestone removes sandy impurities as slag.

Aluminium is extracted by electrolysis of alumina dissolved in molten cryolite; cathode reaction Al3+ + 3e- -> Al.

Copper can occur native; copper is purified by electrolytic refining with impure copper as anode and pure copper as cathode.

An alloy is a mixture of a metal with other elements; alloy layers resist sliding, so the alloy is harder than the pure metal.

Rusting requires both oxygen and water; salt accelerates rusting, and barrier methods or sacrificial protection prevent it.

A more reactive metal displaces a less reactive metal from its salt solution.

Recycling metals saves energy and conserves ore, while extraction can cause environmental damage such as habitat loss and pollution.

Hydrogen is a light diatomic gas H2; a lighted splint gives a pop; it is used in ammonia manufacture and as a fuel.

Oxygen makes up about 21% of air, relights a glowing splint and supports combustion; it is obtained by fractional distillation of liquid air.

Nitrogen makes up about 78% of air and is unreactive because of its strong triple bond.

The Haber process makes ammonia: N2 + 3H2 <=> 2NH3, using an iron catalyst, high pressure and moderate temperature.

Carbon dioxide turns limewater milky and is a greenhouse gas produced by combustion and respiration.

Chlorine is a greenish-yellow gas used for disinfection and bleaching; damp litmus is bleached by chlorine.

Sulfur dioxide is an acidic gas from burning fossil fuels; it causes acid rain and is removed by flue-gas desulfurisation using calcium oxide or calcium carbonate.

Carbon monoxide from incomplete combustion binds to haemoglobin; nitrogen oxides from engines contribute to photochemical smog and acid rain.

Water treatment includes filtration, sedimentation and chlorination; pure water can be checked by its boiling point.

Reaction rate increases with higher concentration or pressure, higher temperature, larger surface area and a catalyst.

Collision theory: particles must collide with at least the activation energy and in the correct orientation for a reaction to occur.

A catalyst lowers activation energy and provides an alternative pathway; it is not used up overall.

An exothermic reaction releases heat and has negative ΔH; an endothermic reaction absorbs heat and has positive ΔH.

Bond breaking is endothermic and bond making is exothermic; ΔH = energy needed to break bonds - energy released forming bonds.

An energy profile shows activation energy from reactants to the peak and ΔH from reactants to products.

At equilibrium the forward and reverse rates are equal and concentrations remain constant, although both reactions continue.

Le Chatelier's principle: increasing pressure shifts equilibrium towards fewer gas moles; increasing temperature shifts it in the endothermic direction.

A catalyst does not change the position of equilibrium; for a given reaction, Kc changes only with temperature.

Organic compounds contain carbon; hydrocarbons contain only carbon and hydrogen.

Alkanes have general formula CnH2n+2, contain only single bonds and are saturated.

Alkenes have general formula CnH2n, contain a C=C double bond and are unsaturated; they decolourise bromine water.

Alcohols contain the -OH functional group; ethanol can be made by fermentation or by hydration of ethene.

Carboxylic acids contain -COOH, are weak acids and react with alcohols to form esters.

Isomers have the same molecular formula but different structural formulae.

A homologous series has the same general formula and functional group, with gradual change in physical properties.

Complete combustion gives carbon dioxide and water; incomplete combustion gives carbon monoxide or carbon.

Addition polymerisation joins alkene monomers; the double bond opens and no other product is formed.

The independent variable is deliberately changed; the dependent variable is measured; control variables are kept constant.

Accuracy is closeness to the true value; precision is closeness of repeated measurements to each other.

Random errors scatter results above and below the true value; systematic errors shift all results in the same direction.

A burette measures variable volumes accurately, a pipette measures one fixed volume, and a measuring cylinder gives approximate volumes.

In titration, read the bottom of the meniscus, use concordant results and identify the endpoint by a colour change.

Gas tests: hydrogen gives a pop, oxygen relights a glowing splint, carbon dioxide turns limewater milky, chlorine bleaches damp litmus and ammonia turns damp red litmus blue.

Flame tests: lithium red, sodium yellow, potassium lilac, calcium orange-red and copper blue-green.

Chromatography can test purity: a pure substance gives one spot, while a mixture gives more than one spot.